Problem 50 - Chemistry An Atoms First Approach

Question

Write electron configurations for a. the cations \(\mathrm{Sr}^{2+}, \mathrm{Cs}^{+}, \mathrm{In}^{+},\) and \(\mathrm{Pb}^{2+}\) b. the anions \(P^{3-}, S^{2-},\) and \(B r^{-}\)

Step-by-Step Solution

Answer

The electron configurations for the given ions are as follows: a. Cations 1. Sr²⁺: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) 2. Cs⁺: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\) 3. In⁺: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10}\) 4. Pb²⁺: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^2 4f^{14} 5d^{10}\) b. Anions 1. P³⁻: \(1s^2 2s^2 2p^6 3s^2 3p^6\) 2. S²⁻: \(1s^2 2s^2 2p^6 3s^2 3p^6\) 3. Br⁻: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\)

Step 1: 1. Sr²⁺ Electron Configuration

Strontium (Sr) has an atomic number of 38. The electron configuration for the neutral Sr atom is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2\] As Sr²⁺ is a cation with a 2+ charge, we will remove two electrons from the 5s orbital. The electron configuration for Sr²⁺ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\]

Step 2: 2. Cs⁺ Electron Configuration

Cesium (Cs) has an atomic number of 55. The electron configuration for the neutral Cs atom is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^1\] As Cs⁺ is a cation with a 1+ charge, we will remove one electron from the 6s orbital. The electron configuration for Cs⁺ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\]

Step 3: 3. In⁺ Electron Configuration

Indium (In) has an atomic number of 49. The electron configuration for the neutral In atom is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^1\] As In⁺ is a cation with a 1+ charge, we will remove one electron from the 5p orbital. The electron configuration for In⁺ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10}\]

Step 4: 4. Pb²⁺ Electron Configuration

Lead (Pb) has an atomic number of 82. The electron configuration for the neutral Pb atom is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^2 4f^{14} 5d^{10} 6p^2\] As Pb²⁺ is a cation with a 2+ charge, we will remove two electrons from the 6p orbital. The electron configuration for Pb²⁺ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^2 4f^{14} 5d^{10}\] b. Electron configurations for anions

Step 5: 1. P³⁻ Electron Configuration

Phosphorus (P) has an atomic number of 15. The electron configuration for the neutral P atom is: \[1s^2 2s^2 2p^6 3s^2 3p^3\] As P³⁻ is an anion with a 3- charge, we will add three electrons in the 3p orbital. The electron configuration for P³⁻ is: \[1s^2 2s^2 2p^6 3s^2 3p^6\]

Step 6: 2. S²⁻ Electron Configuration

Sulfur (S) has an atomic number of 16. The electron configuration for the neutral S atom is: \[1s^2 2s^2 2p^6 3s^2 3p^4\] As S²⁻ is an anion with a 2- charge, we will add two electrons in the 3p orbital. The electron configuration for S²⁻ is: \[1s^2 2s^2 2p^6 3s^2 3p^6\]

Step 7: 3. Br⁻ Electron Configuration

Bromine (Br) has an atomic number of 35. The electron configuration for the neutral Br atom is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^5\] As Br⁻ is an anion with a 1- charge, we will add one electron in the 4p orbital. The electron configuration for Br⁻ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\]

Key Concepts

Understanding Ions

Ions are essential components of chemistry and play a crucial role in chemical reactions and processes. They are atoms or molecules that carry a net electric charge. This charge occurs when the number of protons does not equal the number of electrons in an atom.

• When an atom loses electrons, it becomes positively charged and forms a cation.
• Conversely, when an atom gains electrons, it becomes negatively charged and forms an anion.

This dynamic can lead to different properties and behaviors in ionic compounds. By adjusting electron configurations, atoms can achieve greater stability, typically seeking a full outer shell like the nearest noble gas.

Delving into Cations

Cations are positively charged ions. This positive charge results from losing one or more electrons. In electron configurations, removing these electrons typically starts with the outermost shell, often s or p orbitals. For example:

• Sr²⁺ results from strontium losing two electrons, leaving its electron configuration: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\]
• Cs⁺ arises from cesium losing one electron, simplifying its outer shell to: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\]
• In⁺ results from indium losing an electron, resulting in: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10}\]
• Pb²⁺ is formed when lead loses two electrons, yielding: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^2 4f^{14} 5d^{10}\]

Cations tend to form ionic bonds with anions, resulting in the formation of ionic compounds.

Exploring Anions

Anions are ions with a negative charge, formed when atoms gain extra electrons. This gain typically fills the outermost p or s orbitals closer to achieving a stable electron configuration, resembling that of noble gases.

• P³⁻ sees phosphorus gaining three electrons, altering its configuration to: \[1s^2 2s^2 2p^6 3s^2 3p^6\]
• S²⁻ is generated when sulfur gains two electrons, leading to: \[1s^2 2s^2 2p^6 3s^2 3p^6\]
• Br⁻ forms when bromine gains an additional electron, refining its configuration to: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\]

Gaining electrons helps anions achieve a more stable state, further substantiating the octet rule in chemistry.

Atomic Structure Essentials

Atomic structure is the foundation of understanding ions and electron configurations. Atoms consist of:

• Protons: Positively charged particles located in the nucleus.
• Neutrons: Neutral particles also found in the nucleus.
• Electrons: Negatively charged particles in orbitals around the nucleus.

These components determine an atom’s identity and behavior in chemistry. Electrons arrange themselves in various energy levels or shells, with a typical goal of filling an atom's outermost shell for stability.
The organization of these electrons is depicted through electron configuration, a notation that illustrates which orbitals house the electrons. Mastery of electron configuration, combined with understanding ions, allows us to predict how atoms will interact in compounds and reactions.